Activity series – a list of elements organized according to the ease with which the elements undergo certain chemical reactions
Catalyst – changes the rate of a chemical reaction but can recover unchanged
Chemical equation – represents the identities and relative molecular or molar amounts of the reactants and products in a chemical reaction, using symbols and formulas
Chemical reaction – process by which one or more substances are changed into one or more different substances
Coefficient – a small whole number that appears in front of a formula in a chemical equation
Combustion reaction – occurs when a substance combines with oxygen, releasing a large amount of energy in the form of light and heat
Decomposition reaction – occurs when a single compound undergoes a reaction that produces two or more simpler substances
Double-displacement reaction – occurs when the ions of two compounds exchange places in an aqueous solution to form two new compounds
Electrolysis – the decomposition of a substance by an electric current
Formula equation – represents the reactants and products of a chemical reactions by their symbols or formulas
Precipitate – a solid that is produced as a result of a chemical reaction in solution and that also separates from the solution
Products – the resulting substances following a chemical reaction
Reactants – the original substances in a chemical reaction
Reversible reaction – a chemical reaction in which the products re-form the original reactants
Single-displacement reaction – occurs when one element replaces a similar element in a compound
Synthesis reaction – occurs when two or more substances combine to form a new compound
Describing Chemical Reactions
Signs of Chemical Reactions
A chemical reaction may have occurred if heat and light is released.
Another clue that a chemical reaction has happened is the production of gas bubbles in a substance.
After mixing two solutions, the formation of a precipitate is indicative of the occurence of a chemical reaction.
When a substance changes color, a chemical reaction may have transpired.
Chemical Equations
A chemical equation must identify all reactants and products. It must also include the proper formulas for the reactants and products. Use your knowledge of diatomic molecules (BrINClHOF) and charges of elements to help you with this. Always make sure that the same number of atoms of each element appears on both sides of the equation. Sometimes a coefficient must be used to balance the number of atoms.
Word equations are often a useful step in writing chemical equations. Remember that they only use words to name the reactants and products, but do not give the quantities. Use an → between the different sides of the equation. When reading, the arrow represents “react to yield.”
Next, write a formula equation by substituting the symbols/formulas in for the names of the reactants and products. We place either (g), (s), or (l) to indicate whether the formula is a gas, solid, or a liquid. Note that this equation still does not give the quantities of the reactants or products.
Then, the equation must be balanced by matching the charges and using coefficients. Note: Coefficients show how many molecules of a compound you have.
For instance, the coefficient two in 2H2O indicates that you have two water molecules.
You cannot change subscripts when balancing equations, although it is tempting. Changing subscripts alters the identity of a compound.
For example: 2H2O is NOT the same as H4O. The first shows that there are two separate water molecules; four hydrogens and two oxygens. The latter, which has four hydrogens and one oxygen, does not even exist as a compound.
Always place coefficients in front of chemical formulas. Do not change the subscripts!
Rules to follow when balancing equations:
Balance each different atom separately.
Balance the atoms of elements that are combined, but only occur once on both sides.
Balance polyatomic ions as single units.
Balance hydrogen and oxygen atoms last.
Double check to make sure that the same amount of atoms of each element appears on both sides.
Types of Chemical Reactions
(Also can be found on page 266)
Synthesis Reactions A + X → AX
A common type of this reaction occurs when an element combines with oxygen to form an oxide. Most metals can do this. When this type of reaction happens, Group 2 elements take the formula MO and Group 1 elements take the formula M2O.
For example: 2Mg(s) + O2(g) → 2MgO(s)
When a metal has two different charges, it can form two different oxides.
For example: 2Fe(s) + O2(g) → 2FeO(s)
This reaction can also occur between Group 1 and 2 elements with sulfur to produce sulfides. Group 2 elements take the formula MS and Group 1 elements take the formula M2S.
Nonmetals can also use synthesis reactions with oxygen to form oxides.
For example: C(s) + O2(g) → CO2(g)
When Group 1 metals react with halogens, ionic compounds are formed with the formula MX (M is the metal; X is the halogen).
For example: 2Na(s) + Cl2(g) → 2NaCl(s)
Group 2 metals and halogens react to take the formula MX2. Because of its reactivity, fluorine combines with most metals.
For example: Mg(s) + F2(g) → MgF2(s)
When an oxide of an active metal reacts with water, a metal hydroxide is created.
For example: CaO(s) + H2O(l) → Ca(OH)2(s)
Most nonmetal oxides react with water to produce oxyacids.
For example: SO2(g) + H2O(l) → H2SO3(aq)
When some nonmetal oxide reacts with metal oxides, it forms a salt.
For example: CaO(s) + SO2(g) → CaSO3(s)
Decomposition Reactions AX → A + X
These reactions can be simple when they involve turning binary compounds into their original elements. In most cases, a decomposition reaction requires electricity or heat. When done with electricity, it is called electrolysis. This method is used when water is decomposed into hydrogen and oxygen.
For example: 2H2O(l) electricity→ 2H2(g) + O2(g)
Oxides decompose using heat. Oxygen was actually discovered through this process in the 1700s. Mercury(II) oxide, when heated, forms mercury and oxygen.
For example: 2HgO(s) Δ→ 2Hg(l) + O2(g)
When heated, metal carbonates form metal oxide and carbon dioxide gas.
For example: CaCO3(s) Δ→ CaO(s) + CO2(g)
Metal hydroxides, excluding those with metals in Group I, when heated, form a metal oxide and water.
For example: Ca(OH)2(s) Δ→ CaO(s) + H2O(g)
When heated and influenced by a catalyst, metal chlorates decompose to form a metal chloride and oxygen.
For example: 2KClO3(s) ΔMnO2(s)→ 2KCl(s) + 3O2(g)
Nonmetal oxides and water are produced when some acids decompose.
For example: H2CO3(aq) → CO2(g) + H2O(l)
Single-Displacement Reactions
This type of reaction often occurs in aqueous solutions. The formulas are A + BX → AX + B.
In a compound, a metal can displace another metal.
For example: 2Al(s) + 3Pb(NO3)2(aq) → 3Pb(s) + 2Al(NO3)3(aq)
A metal can also displace hydrogen in water.
For example: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
Hydrogen found in an acid can be displaced by a metal as well.
For example: Mg(s) + 2HCl(aq) → H2(g) + MgCl2(aq)
Halogens can displace each other.
For example: Cl2(g) + 2KBr(aq) → 2KCl(aq) + Br2(l)
Double-Displacement Reactions AX + BY → AY + BX
This reaction usually forms a precipitate, a gas, or water.
Precipitates arise through the combination of cations and anions during the formation of an insoluble compound.
For example: 2KI(aq) + Pb(NO3)2(aq) → PbI2(s) + 2KNO3(aq)
Gas forms as a product of the reaction. It bubbles out of the mixture.
For example: FeS(s) + 2HCl(aq) → H2S(g) + FeCl2(aq)
Water can form as a product as well.
For example: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Combustion Reactions
Combustion reactions occur when a substance combines with oxygen and releases energy.
(Also can be found on page 286)
For example: 2H2(g) + O2(g) → 2H2O(g)
Activity Series of the Elements
When a metal has a larger activity, it loses electrons more easily. When a nonmetal has a larger activity, it gains electrons more readily.
The series is listed from the most-active element to the least-active element.
Single-displacement reactions help to decipher the order of an element’s activity.
An element can replace any element below it in the series, but no element above it.
Group 17 elements, the halogens, are actually listed properly according to their activity on the periodic table.
It also helps to interpret if a written chemical equation can actually occur. (See if one element can or cannot replace the other based on activity)
For example: 2Al(s) + 3ZnCl2(aq) → 3Zn(s) + 2AlCl3(aq)
Works Cited
Holt, Rinehart and Winston, ed. Modern Chemistry. Austin: Harcourt Education Company, 2006.
Chemical Equations and Reactions
This chapter discusses the following topics:
Describing Chemical Reactions
Types of Chemical Reactions
Activity Series of the Elements
Key Terms
Describing Chemical Reactions
Signs of Chemical Reactions
Chemical Equations
For instance, the coefficient two in 2H2O indicates that you have two water molecules.
Rules to follow when balancing equations:
Types of Chemical Reactions
Synthesis Reactions A + X → AX
Decomposition Reactions AX → A + X
Single-Displacement Reactions
Double-Displacement Reactions AX + BY → AY + BX
Combustion Reactions
Activity Series of the Elements
Works Cited
Holt, Rinehart and Winston, ed. Modern Chemistry. Austin: Harcourt Education Company, 2006.